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12.2 GROUP IIA: ALKALINE EARTHS

Some properties of the alkaline earths have already been described in Sec. 4.1. You should review that section now. The last member of the group, Ra, is radioactive and will not be considered here. All alkaline earths are silvery-gray metals which are ductile and relatively soft. However, Table 12.2 shows that they are much denser than the group IA metals, and their melting points are significantly higher. They are also harder than the alkali metals. This may be attributed to the general valence electron configuration ns² for the alkaline earths, which involves two electrons per metal atom in metallic bonding (instead of just one as in an alkali metal).

First and second ionization energies for the alkaline earths (corresponding to removal of the first and second valence electrons) are relatively small, but the disruption of an octet by removal of a third electron is far more difficult (see Table 12.2). Like the alkali metals, the alkaline-earth atoms lose electrons easily, and so they are good reducing agents. Except for Be (which has the largest ionization energy), they react directly with most nonmetallic elements, forming hydrides: MH₂ (M = Mg, Ca, Sr, Ba); halides: MF₂, MCl₂, MBr₂, and MI₂; oxides: MO; and sulfides: MS. In all these compounds the alkaline-earth elements occur as dipositive ions, Mg²⁺, Ca²⁺, Sr²⁺, or Ba²⁺.

Similar compounds of Be can be formed by roundabout means, but not by direct combination of the elements. Moreover, the Be compounds are

TABLE 12.2 Properties of the Group IIA Alkaline-Earth Metals.

More covalent than ionic. The Ba²⁺ ion has a very small radius (31 pm) and is therefore capable of distorting (polarizing) the electron cloud of an anion in its vicinity. Therefore all bonds involving Be have considerable covalent character, and the chemistry of Be is significantly different from that of the other members of group IIA.

Other trends among the data in Table 12.2 are what we would expect. Ionization energies and electronegativities decrease from top to bottom of the group, and atomic and ionic radii increase. The radii of +2 alkaline-earth ions are much smaller than the +1 alkali-metal ions of the same period (compare Tables 12.1 and 12.2), because the greater nuclear charge holds the inner shells more tightly. This effect is sufficiently large that an alkaline earth below and to the right of a given alkali metal in the periodic table often has nearly the same ionic radius. Thus Na⁺ (95 pm), can fit into exactly the same type of crystal lattice as Ca²⁺ (99 pm), and these two elements are often found in the same minerals. The same is true of K⁺ and Ba²⁺.

Similarity of ionic radii also leads to related properties for Li and Mg. Since these two elements are adjacent along a diagonal line from the upper left to the lower right in the periodic table, their similarity is called a diagonal relationship. Diagonal relationships are mainly evident in the second and third periods: Be is similar to Al, and B is like Si in many ways.

Farther toward the right-hand side of the table such relationships are less pronounced. The most striking similarity between Li and Mg is their ability to form covalent bonds with elements of average electronegativity, such as C, while forming fairly ionic compounds with more electronegative elements, such as O or F. Two examples of covalent compounds are ethyllithium, CH₃CH₂Li, and diethylmagnesium, (CH₃CH₂)₂Mg. Such compounds are likely in the case of Li and Mg but not the alkali or alkaline earths below them, because Li⁺ and Mg²⁺ are small enough to be strongly polarizing and thus form bonds with considerable covalent character.

Chemical Reactions and Compounds

As in the case of the alkali metals, the most important and abundant alkaline earths, Mg and Ca, are in the third and fourth periods. Be is rare, although its strength and low density make it useful in certain special

alloys. Sr and Ba occur naturally as the relatively insoluble sulfates SrSO₄ (strontianite) and BaSO₄ (barite), but these two elements are of minor commercial importance.

The most common ores of Mg and Ca are dolomite, MgCO₃•CaCO₃, after which an entire mountain range in Italy is named, and limestone, CaCO₃, an important building material. Mg is also recovered from seawater on a wide scale. The oxides of the alkaline earths are commonly obtained by heating the carbonates. For example, lime, CaO, is obtained from limestone as follows:

CaCO₃(s) CaO(s) + CO₂(g)

Except for BeO, which is covalently bonded, alkaline-earth oxides contain O^2– ions and are strongly basic. When treated with water (a process known as slaking), they are converted to hydroxides:

CaO(s) + H₂O(l) → Ca(OH)₂(s)

Ca(OH)₂ (slaked lime) is an important strong base for industrial applications, because it is cheaper than NaOH.

MgO has an extremely high melting point (2800°C) because of the close approach and large charges of its constituent Mg²⁺ and O^2– ions in the crystal lattice. As a solid it is a good electrical insulator, and so it is used to surround metal-resistance heating wires in electric ranges. MgO is also used to line high-temperature furnaces. When converted to the hydroxide, Mg finds a different use. Mg(OH)₂ is quite insoluble in water, and so it does not produce a high enough concentration of hydroxide ions to be caustic. It is basic, however, and gram for gram can neutralize nearly twice the quantity of acid that NaOH can. Consequently a suspension of Mg(OH)₂ water (milk of magnesia) makes an excellent antacid, for those who can stand its taste.

Because the carbonate ion behaves as a Brönstedt-Lowry base, carbonate salts dissolve in acidic solutions. In nature, water often becomes acidic because the acidic oxide CO₂ is present in the atmosphere. When CO₂ from the air dissolves in water, it can help dissolve limestone:

CO₂(g) + H₂O(l) + CaCO₃(s) Ca²⁺(aq) + HCO₃^–(aq)

This reaction often occurs underground as rainwater saturated with CO₂

TABLE 12.3 Properties of the Group IIIA Elements.

seeps through a layer of limestone. Caves from which the limestone has been dissolved are often prevalent in areas where there are large deposits of CaCO₃. In addition, the groundwater and well water in such areas becomes hard. Hard water contains appreciable concentrations of Ca²⁺, Mg²⁺ , and certain other metal ions. These form insoluble compounds with soap, causing curdy, scummy precipitates. Hard water can be softened by adding Na₂CO₃, washing soda, which precipitates CaCO₃, or by ion exchange, a process in which the undesirable Ca²⁺ and Mg²⁺ ions are replaced in solution by Na⁺ ions, which do not precipitate soap. Most home water softeners work on the latter principle.

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